"CHEMISTRY REVIEW" PROJECT PAGES
This Project Page first appeared in the September 1996 issue of Chemistry Review, Volume 6,
Number 1, Pages 26 and 27. Chemistry Review is published four times during the academic year by
Philip Allan Updates and is a journal for post-16 students. It contains a variety of interesting and
colourful articles aimed at 16-19 year-olds taking mainly AS and A2 courses in chemistry.
NOTE: Project Page is designed to help you think about your investigation. It is not intended to be a
set of instructions for practical work and does not include a list of safety precautions. CHEMISTRY
REVIEW accepts no responsibility if Project Page is used in any way as a set of instructions.
Reactions that don’t seem to add up
Before you began to read this section you may well have assumed that chemical reactions are
commutative. That is, just as you expect to get the same answer when you add 2 to 3 as you get
when you add 3 to 2, so you might expect get the same product formed when you add substance A to
substance B as you do by adding substance B to substance A. In some cases, however, this just
does not happen. The outcome of the chemical reaction depends very definitely on the order in which
the substances are mixed together. Such systems offer the possibility of some simple but intriguing
projects.
Copper(II) sulphate – sodium hydroxide reaction
The reaction between copper(Il) sulphate and sodium hydroxide solutions is a good place to start. If
you slowly add one to the other while stirring, you will get a precipitate of copper(II) hydroxide,
Cu(OH)2. If you mix the same solutions the other way round, you will again get a precipitate, but this
time it will contain sulphate ions as well as copper(Il) and hydroxide ions.
You can find out some information about the formulas the precipitates while they are being formed by
using a pH meter and/or a conductivity meter. In both cases you will need to take readings from the
meters as you add one solution to a fixed volume of the other one. Do make sure that you are using a
sufficient volume to cover the lower part of the pH or conductivity probe. If you plot your results on a
graph, you will see a very marked change when the reaction is complete (see Figures 1 and 2). You
can use the amounts of the reactants needed to achieve this point to find the ratio of copper(Il) and
hydroxide ions in the two precipitates.
pH
End point
Volume of added sodium hydroxide solution
Figure 1 Typical graph of pH against volume of
sodium hydroxide solution added to a fixed
volume of copper(II) sulphate solution.
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Conductivity
End point
Volume of added sodium hydroxide solution
Figure 2 Typical graph of conductivity against volume of sodium hydroxide
solution added to a fixed volume of copper(II) sulphate solution
You can find but more about the formula of the precipitate by first filtering it off using a Buchner funnel
and drying it (at room temperature to prevent decomposition to copper oxide). Then you can
determine how much copper and how much sulphate is present. You can then calculate the amount
hydroxide by difference.
There are two approaches you might use to finding out ho much copper the precipitate contains. One
way is to add weighed sample of dried precipitate to dilute ethanoic acid so that the pH is low enough
to redissolve the precipitate of Cu(OH)2. Excess potassium iodide solution is added which reduces
the Cu(II) to Cu(I):
2Cu2+ + 2I– →
I2 + 2Cu+
The liberated iodine is then titrated with sodium thiosulphate solution using starch indicator.
Alternatively, you could add your weighed sample to ammonia solution. All the copper ions
present complex with ammonia molecules and the product is deep blue in colour:
Cu2+ + 4NH3
→
[Cu(NH3)4]2+
The intensity of the colour is proportional to the concentration of the complex, and therefore to the
amount of copper ions present in the original sample of precipitate. You need to use a colorimeter to
compare the intensity of the colour obtained using a sample of your precipitate with that obtained from
standard samples of copper ions added to ammonia solution.
You can easily find out whether there are measurable amounts of sulphate ions in your precipitate by
adding it to dilute hydrochloric acid and then adding a solution of barium chloride. The sulphate will be
precipitated as white insoluble barium sulphate. Notice that the presence of hydrochloric acid is
essential to ensure that the precipitate consists entirely of barium sulphate and that there is no barium
hydroxide. The exact amount of sulphate is best found by weighing (what chemists call a gravimetric
method). To do this, add a weighed amount of dry precipitate to hydrochloric acid and precipitate the
barium sulphate with an excess of barium chloride. Filter off the precipitate (through a sintered glass
crucible if you have access to one) and dry in an oven at just over 100 °C until its mass stops
changing.
By combining all the information you have gathered from your pH and conductivity measurements and
from your analysis experiments you should be able to decide on the composition of the precipitates
formed by mixing your solutions in a different order.
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Extending your investigations
You can use the model of copper(II) sulphate to investigate the ways in which a range of other
solutions react with sodium hydroxide. What happens if you use copper(II) nitrate or chloride instead
of the sulphate? What happens if you use the sulphates of other metals such as zinc, nickel,
magnesium or manganese? pH and conductivity measurements can be used in all of these systems.
(You will need to look up methods for analysing other ions to provide extra information.) It will become
evident that not all reactions add up in the same way as 2 plus 3 and 3 plus 2!
Do the formulas of crystals always add up?
When hot, saturated solutions of two different metal sulphates are mixed and cooled only one kind of
crystal separates out. These crystals contain both kinds of metal ions, and finding out their exact
composition and shape can provide yet more starting points for projects.
Some pairs of sulphates produce crystals of which the composition depends on the proportions in
which the saturated solutions were mixed. Not only can two salts of this kind form mixed crystals but,
if you add a crystal of either sulphate to a saturated solution of the other one, the crystal will continue
to grow, the second salt being deposited on top of the first one. This is an example of what is called
an overgrowth.
When other pairs of sulphates are mixed the outcome is quite different; the crystals formed always
have exactly the same composition, whatever the composition of the solution from which they
crystallise. Such crystals are called double salts. One group is referred to as 'alums'; these have the
general formula A2SO4. B2(SO4)3. 24H2O. The other group has the general formula A2SO4. BSO4.
6H2O, where A and B represent the symbols of different cations.
A number of questions await your investigation. In what ways do specific pairs of sulphates behave
on mixing? In what ways are the shapes of the crystals formed related to the shapes of the original
sulphate crystals? How is the number of molecules of water of crystallisation in the parent sulphates
related to the number in mixed crystals? What happens if you mix solutions made from pairs of
different double salts?
Think about how you can use ways of analysing for metal ions to help you decide on the crystal
composition. A micro- scope will help you to see the differences in crystal habit and this may
introduce a new and quite striking aspect to your chemical studies. The choice of sulphates you can
investigate is wide. Sulphates of magnesium, iron(II), copper(II), zinc, ammonium, potassium, iron(III),
chromium(III) and aluminium will provide many combinations worth investigating. You will find that the
formulas of mixed sulphates are not exactly crystal clear.
Derek Denby
Derek Denby is Head of Chemistry at John Leggott College, Scunthorpe.
The original article was written by Derek Denby. We are grateful to Derek for allowing us to
reproduce it here.
This page is free for your personal use, but the copyright remains with Philip Allan Updates. Please do
not copy it or disseminate it in any way.
Chemistry Review is indebted to Don Ainley, who has helped to prepare this article for the Web.
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